Standard electrode reduction potentials

TABLE 4.2 Some Standard Electrode (Reduction) Potentials in Aqueous Solution at 25°C... 

Table 6.2 Selected Standard Electrode Reduction Potentials E d in Water (V vs. NHE) at 298.15 K, Assuming Unimolar Concentrations or Activities for Solutes and Unit Pressures or Fugacities for Gases (Ordered First by Potential, then Alphabetically) ...

In this book standard electrode (reduction) potentials will be employed, as recommended by the I.U.P.A.C. Table 8.1 f gives a list of such potentials. The reactions are written as reduction processes, e.g., Cu + 2e" -> Cu, A table of standard oxidation potentials would simply reverse all of the signs and the corresponding reactions would be oxidations, e.g., Cu Cu " + 26 ,... 

The standard electrode potential at the above mentioned standard state conditions is denoted by °. For the MCAT, the values of the standard electrode (reduction) potentials will be given to you if you are required to solve such a question. Do not try to memorize those values. The standard electrode potentials are based on an arbitration with reference to standard hydrogen electrode. The standard hydrogen electrode potential is considered to be 0 volt. 

The potential of these half-cell reactions can be determined by comparison with the hydrogen electrode all under standard conditions, which is unit concentration for ions in solution and 1 atm pressure for gases at 25°C. The value of these standard electrode reduction potentials is given in Table 9.4. The standard cell potential of reaction (9.15) is... 

In Appendix D, we have modified the table of Standard Electrode (Reduction) Potentials at 25 °C so that it now includes a column with the cell notation for the half-reactions. 

When a biochemical half-reaction involves the production or consumption of hydrogen ions, the electrode potential depends on the pH. When reactants are weak acids or bases, the pH dependence may be complicated, but this dependence can be calculated if the pKs of both the oxidized and reduced reactants are known. Standard apparent reduction potentials E ° have been determined for a number of oxidation-reduction reactions of biochemical interest at various pH values, but the E ° values for many more biochemical reactions can be calculated from ArG ° values of reactants from the measured apparent equilibrium constants K. Some biochemical redox reactions can be studied potentiometrically, but often reversibility cannot be obtained. Therefore a great deal of the information on reduction potentials in this chapter has come from measurements of apparent equilibrium constants. 

Chlorostannate and chloroferrate [110] systems have been characterized but these metals are of little use for electrodeposition and hence no concerted studies have been made of their electrochemical properties. The electrochemical windows of the Lewis acidic mixtures of FeCh and SnCh have been characterized with ChCl (both in a 2 1 molar ratio) and it was found that the potential windows were similar to those predicted from the standard aqueous reduction potentials [110]. The ferric chloride system was studied by Katayama et al. for battery application [111], The redox reaction between divalent and trivalent iron species in binary and ternary molten salt systems consisting of 1-ethyl-3-methylimidazolium chloride ([EMIMJC1) with iron chlorides, FeCb and FeCl j, was investigated as possible half-cell reactions for novel rechargeable redox batteries. A reversible one-electron redox reaction was observed on a platinum electrode at 130 °C. 

Determination of Standard Oxidation-Reduction Potentials.—In principle, the determination of the standard potential of an oxidation-reduction system involves setting up electrodes containing the oxidized and reduced states at known activities and measuring the potential B by combination with a suitable reference electrode insertion of the value of B in the appropriate form of equation (3) then permits B to be calculated. The inert metal employed in the oxidation-reduction electrode is frequently of smooth platinum, clthough platinized platinum, mercury and particularly gold are often used. 

When the pH is specified, each biochemical half reaction makes an independent contribution to the apparent equilibrium constant K for the reaction written in terms of reactants rather than species. The studies of electochemical cells have played an important role in the development of biochemical thermodynamics, as indicated by the outstanding studies by W. Mansfield Clarke (1). The main source of tables of ° values for biochemical half reactions has been those of Segel (2). Although standard apparent reduction potentials ° can be measured for some half reactions of biochemical interest, their direct determination is usually not feasible because of the lack of reversibility of the electrode reactions. However, standard apparent reduction potentials can be calculated from for oxidoreductase reactions. Goldberg and coworkers (3) have compiled and evaluated the experimental determinations of apparent equilibrium constants and standard transformed enthalpies of oxidoreductase reactions, and their tables have made it possible to calculate ° values for about 60 half reactions as functions of pH and ionic strength at 298.15 K (4-8). 

Table 2 Standard oxidation-reduction potentials 25° C, volts Di. hydrogen electrode...

Each electrode reaction, anode and cathode, or half-cell reaction has an associated energy level or electrical potential (volts) associated with it. Values of the standard equilibrium electrode reduction potentials E° at unit activity and 25°C may be obtained from the literature (de Bethune and Swendeman Loud, Encyclopedia of Electrochemistry, Van Nostrand Reinhold, 1964). The overall electrochemical cell equilibrium potential either can be obtained from AG values or is equal to the cathode half-cell potential minus the anode half-cell potential, as shown above. 

Measurement of Standard Single Electrode Reduction Potentials... 

The single electrode reduction potentials are measured with respect to the SHE. The SHE, which is based on the reaction, 2H+ + 2e -o- H2 on platinized platinum, is assigned the reduction potential of zero when the activity of the hydrogen ion is unity and the pressure of H2 is 1 atm. Figure 4.1.6 illustrates the schematic for measuring the standard single reduction potential of Cu /Cu and Zn /Zn systems. 

The Standard Oxidation-Reduction Potential Eq of Some Biological Compounds at pH 7.0 Referred to the Standard Hydrogen Electrode... 

To measure Eo, it is only necessary to measure E when Ox) = Red). Eo is termed the standard oxidation-reduction potential, it is the potential which is measured against the hydrogen electrode at the pH of a normal solution of HCl (pH == 0, since log 1 = 0). It has become customary to state the potential at, or around, pH 7-0 and this standard potential is designated by the symbol Eo-... 

Finally, Table 2.2 presents data for the standard oxidation-reduction potentials EJ (equation (2.149)) of some metal electrodes, which are most frequently used as substrates for studies of electrochemical phase formation. 

V. 19-8. In a calomel electrode, reduction potential depends on the chloride potential. Therefore, the standard reduction potential for a calomel electrode has a different chloride concentration from the saturated calomel electrode. 19-9. Dry cells and lead-acid cells "run down" as the concentrations of reactants and products eventually reach their equilibrium values, where A,G and Eceii both become 0. This does not happen in a fuel cell because fuel is continuously added. 19-10. Both AI and Zn can be used because they are more active than Fe ... 

For example, for iron in aqueous electrolytes, tlie tliennodynamic warning of tlie likelihood of corrosion is given by comparing tlie standard electrode potential of tlie metal oxidation, witli tlie potential of possible reduction reactions. 

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